Gravimetric analysisGravimetric analysis describes a set of methods used in analytical
chemistry for the quantitative determination of an analyte (the ion
being analyzed) based on its mass. The principle behind this type of
analysis is that once an ion's mass has been determined as a unique
compound, that known measurement can then be used to determine the
same analyte's mass in a mixture, as long as the relative quantities
of the other constituents are known.[1]
The four main types of this method of analysis are precipitation,
volitilization, electro-analytical and miscellaneous physical
method.[2] The methods involve changing the phase of the analyte in
order to separate it in its pure form from the original mixture and
are quantitative measurements.

Precipitation method[edit]
The precipitation method is the one used for the determination of the
amount of calcium in water. Using this method, an excess of oxalic
acid, H2C2O4, is added to a measured, known volume of water. By adding
a reagent, here ammonia, the calcium will precipitate as calcium
oxalate. The proper reagent, when added to aqueous solution, will
produce highly insoluble precipitates from the positive and negative
ions that would otherwise be soluble with their counterparts (equation
1).[3]
The reaction is:
Formation of calcium oxalate:
Ca2+(aq) + C2O42 -→ CaC2O4
The precipitate is collected, dried and ignited to high (red) heat
which converts it entirely to calcium oxide.
The reaction is pure calcium oxide formed
CaC2O4 → CaO(s) + CO(g)+ CO2(g)
The pure precipitate is cooled, then measured by weighing, and the
difference in weights before and after reveals the mass of analyte
lost, in this case calcium oxide.[4][5] That number can then be used
to calculate the amount, or the percent concentration, of it in the
original mix.[2][4][5]
Types of volatilization methods[edit]
In volatilization methods, removal of the analyte involves separation
by heating or chemically decomposing a volatile sample at a suitable
temperature.[2][6] In other words, thermal or chemical energy is used
to precipitate a volatile species.[7] For example, to determine the
water content of a compound by vaporizing the water using thermal
energy (heat). Heat can also be used, if oxygen is present, for
combustion to isolate the suspect species and obtain the desired
results.
The two most common gravimetric methods using volatilization are those
for water and carbon dioxide.[2] An example of this method is the
isolation of sodium hydrogen bicarbonate (the main ingredient in most
antacid tablets) from a mixture of carbonate and bicarbonate.[2] The
total amount of this analyte, in whatever form, is obtained by
addition of an excess of dilute sulfuric acid to the analyte in
solution.
In this reaction, nitrogen gas is introduced through a tube into the
flask which contains the solution. As it passes through, it gently
bubbles. The gas then exits, first passing a drying agent (here CaSO4,
the common desiccant Drierite). It then passes a mixture of the drying
agent and sodium hydroxide which lays on asbestos or Ascarite II, a
non-fibrous silicate containing sodium hydroxide[8] The mass of the
carbon dioxide is obtained by measuring the increase in mass of this
absorbent.[2] This is performed by measuring the difference in weight
of the tube in which the ascarite contained before and after the
procedure.
The calcium sulfate (CaSO4) in the tube retains carbon dioxide
selectively as it's heated, and thereby, removed from the solution.
The drying agent absorbs any aerosolized water and/or water vapor
(reaction 3.). The mix of the drying agent and NaOH absorbs the CO2
and any water that may have been produced as a result of the
absorption of the NaOH (reaction 4.).[9]
The reactions are:
Reaction 3 - absorption of water
NaHCO3(aq) + H2SO4(aq) → CO2(g) + H2O(l) + NaHSO4(aq).[9]
Reaction 4. Absorption of CO2 and residual water
CO2(g) + 2 NaOH(s) → Na2CO3(s) + H2O(l).[9]
Volatilization methods[edit]
Volatilization methods can be either direct or indirect. Water
eliminated in a quantitative manner from many inorganic substances by
ignition is an example of a direct determination. It is collected on a
solid desiccant and its mass determined by the gain in mass of the
desiccant.
Another direct volatilization method involves carbonates which
generally decompose to release carbon dioxide when acids are used.
Because carbon dioxide is easily evolved when heat is applied, its
mass is directly established by the measured increase in the mass of
the absorbent solid used.[10][11]
Determination of the amount of water by measuring the loss in mass of
the sample during heating is an example of an indirect method. It is
well known that changes in mass occur due to decomposition of many
substances when heat is applied, regardless of the presence or absence
of water. Because one must make the assumption that water was the only
component lost, this method is less satisfactory than direct methods.
This often fault and misleading assumption has proven to be wrong on
more than a few occasions. There are many substances other than water
loss that can lead to loss of mass with the addition of heat, as well
as a number of other factors that may contribute to it. The widened
margin of error created by this all-too-often false assumption is not
one to be lightly disregarded as the consequences could be
far-reaching.
Nevertheless, the indirect method, although less reliable than direct,
is still widely used in commerce. For example, it's used to measure
the moisture content of cereals, where a number of imprecise and
inaccurate instruments are available for this purpose.

Procedure[edit]

The sample is dissolved, if it is not already in solution.
The solution may be treated to adjust the pH (so that the proper
precipitate is formed, or to suppress the formation of other
precipitates). If it is known that species are present which interfere
(by also forming precipitates under the same conditions as the
analyte), the sample might require treatment with a different reagent
to remove these interferents.
The precipitating reagent is added at a concentration that favors the
formation of a "good" precipitate (see below). This may require low
concentration, extensive heating (often described as "digestion"), or
careful control of the pH. Digestion can help reduce the amount of
coprecipitation.
After the precipitate has formed and been allowed to "digest", the
solution is carefully filtered. The filter is used to collect the
precipitate; smaller particles are more difficult to filter.

Depending on the procedure followed, the filter might be a piece of
ashless filter paper in a fluted funnel, or a filter crucible. Filter
paper is convenient because it does not typically require cleaning
before use; however, filter paper can be chemically attacked by some
solutions (such as concentrated acid or base), and may tear during the
filtration of large volumes of solution.
The alternative is a crucible whose bottom is made of some porous
material, such as sintered glass, porcelain or sometimes metal. These
are chemically inert and mechanically stable, even at elevated
temperatures. However, they must be carefully cleaned to minimize
contamination or carryover(cross-contamination). Crucibles are often
used with a mat of glass or asbestos fibers to trap small particles.
After the solution has been filtered, it should be tested to make sure
that the analyte has been completely precipitated. This is easily done
by adding a few drops of the precipitating reagent; if a precipitate
is observed, the precipitation is incomplete.

After filtration, the precipitate – including the filter paper or
crucible – is heated, or charred. This accomplishes the following:

The remaining moisture is removed (drying).
Secondly, the precipitate is converted to a more chemically stable
form. For instance, calcium ion might be precipitated using oxalate
ion, to produce calcium oxalate (CaC2O4); it might then be heated to
convert it into the oxide (CaO). It is vital that the empirical
formula of the weighed precipitate be known, and that the precipitate
be pure; if two forms are present, the results will be inaccurate.
The precipitate cannot be weighed with the necessary accuracy in place
on the filter paper; nor can the precipitate be completely removed
from the filter paper in order to weigh it. The precipitate can be
carefully heated in a crucible until the filter paper has burned away;
this leaves only the precipitate. (As the name suggests, "ashless"
paper is used so that the precipitate is not contaminated with ash.)

After the precipitate is allowed to cool (preferably in a desiccator
to keep it from absorbing moisture), it is weighed (in the crucible).
To calculate the final mass of the analyte, the starting mass of the
empty crucible is subtracted from the final mass of the crucible
containing the sample. Since the composition of the precipitate is
known, it is simple to calculate the mass of analyte in the original
sample.

Example[edit]
A chunk of ore is to be analyzed for sulfur content. It is treated
with concentrated nitric acid and potassium chlorate to convert all of
the sulfur to sulfate (SO2−
4). The nitrate and chlorate are removed by treating the solution with
concentrated HCl. The sulfate is precipitated with barium (Ba2+) and
weighed as BaSO4.
Advantages[edit]
Gravimetric analysis, if methods are followed carefully, provides for
exceedingly precise analysis. In fact, gravimetric analysis was used
to determine the atomic masses of many elements to six figure
accuracy.
GravimetryGravimetry provides very little room for instrumental error
and does not require a series of standards for calculation of an
unknown. Also, methods often do not require expensive equipment.
Gravimetric analysis, due to its high degree of accuracy, when
performed correctly, can also be used to calibrate other instruments
in lieu of reference standards.
Disadvantages[edit]
Gravimetric analysisGravimetric analysis usually only provides for the analysis of a
single element, or a limited group of elements, at a time. Comparing
modern dynamic flash combustion coupled with gas chromatography with
traditional combustion analysis will show that the former is both
faster and allows for simultaneous determination of multiple elements
while traditional determination allowed only for the determination of
carbon and hydrogen. Methods are often convoluted and a slight
mis-step in a procedure can often mean disaster for the analysis
(colloid formation in precipitation gravimetry, for example). Compare
this with hardy methods such as spectrophotometry and one will find
that analysis by these methods is much more efficient.
Steps in a gravimetric analysis[edit]
After appropriate dissolution of the sample the following steps should
be followed for successful gravimetric procedure:
1. Preparation of the Solution: This may involve several steps
including adjustment of the pH of the solution in order for the
precipitate to occur quantitatively and get a precipitate of desired
properties, removing interferences, adjusting the volume of the sample
to suit the amount of precipitating agent to be added.
2. Precipitation: This requires addition of a precipitating agent
solution to the sample solution. Upon addition of the first drops of
the precipitating agent, supersaturation occurs, then nucleation
starts to occur where every few molecules of precipitate aggregate
together forming a nucleous. At this point, addition of extra
precipitating agent will either form new nuclei or will build up on
existing nuclei to give a precipitate. This can be predicted by Von
Weimarn ratio where, according to this relation the particle size is
inversely proportional to a quantity called the relative
supersaturation where
Relative supersaturation = (Q – S)/S
The Q is the concentration of reactants before precipitation, S is the
solubility of precipitate in the medium from which it is being
precipitated. Therefore, in order to get particle growth instead of
further nucleation we need to make the relative supersaturation ratio
as small as possible. The optimum conditions for precipitation which
make the supersaturation low are:
a. Precipitation using dilute solutions to decrease Q b. Slow addition
of precipitating agent to keep Q as low as possible c. Stirring the
solution during addition of precipitating agent to avoid concentration
sites and keep Q low d. Increase solubility by precipitation from hot
solution e. Adjust the pH in order to increase S but not a too much
increase np as we do not want to lose precipitate by dissolution f.
Usually add a little excess of the precipitating agent for
quantitative precipitation and check for completeness of the
precipitation
3. Digestion of the precipitate: The precipitate is left hot (below
boiling) for 30 min to 1 hour in order for the particles to be
digested. Digestion involves dissolution of small particles and
reprecipitation on larger ones resulting in particle growth and better
precipitate characteristics. This process is called Ostwald ripening.
An important advantage of digestion is observed for colloidal
precipitates where large amounts of adsorbed ions cover the huge area
of the precipitate. Digestion forces the small colloidal particles to
agglomerate which decreases their surface area and thus adsorption.
You should know that adsorption is a major problem in gravimetry in
case of colloidal precipitate since a precipitate tends to adsorb its
own ions present in excess, Therefore, forming what is called a
primary ion layer which attracts ions from solution forming a
secondary or counter ion layer. Individual particles repel each other
keeping the colloidal properties of the precipitate. Particle
coagulation can be forced by either digestion or addition of a high
concentration of a diverse ions strong electrolytic solution in order
to shield the charges on colloidal particles and force agglomeration.
Usually, coagulated particles return to the colloidal state if washed
with water, a process called peptization.
4. Washing and Filtering the Precipitate: It is crucial to wash the
precipitate very well in order to remove all adsorbed species which
will add to weight of precipitate. One should be careful nor to use
too much water since part of the precipitate may be lost. Also, in
case of colloidal precipitates we should not use water as a washing
solution since peptization would occur. In such situations dilute
nitric acid, ammonium nitrate, or dilute acetic acid may be used.
Usually, it is a good practice to check for the presence of
precipitating agent in the filtrate of the final washing solution. The
presence of precipitating agent means that extra washing is required.
FiltrationFiltration should be done in appropriate sized Gooch or ignition
filter paper.
5. Drying and Ignition: The purpose of drying (heating at about
120-150 oC in an oven) or ignition in a muffle furnace at temperatures
ranging from 600-1200 oC is to get a material with exactly known
chemical structure so that the amount of analyte can be accurately
determined.
6. Precipitation from Homogeneous Solution: In order to make Q minimum
we can, in some situations, generate the precipitating agent in the
precipitation medium rather than adding it. For example, in order to
precipitate iron as the hydroxide, we dissolve urea in the sample.
Heating of the solution generates hydroxide ions from the hydrolysis
of urea. Hydroxide ions are generated at all points in solution and
thus there are no sites of concentration. We can also adjust the rate
of urea hydrolysis and thus control the hydroxide generation rate.
This type of procedure can be very advantageous in case of colloidal
precipitates.
Solubility in the presence of diverse ions[edit]
As expected from previous information, diverse ions have a screening
effect on dissociated ions which leads to extra dissociation.
Solubility will show a clear increase in presence of diverse ions as
the solubility product will increase. Look at the following example:
Find the solubility of AgCl (Ksp = 1.0 x 10−10) in 0.1 M NaNO3. The
activity coefficients for silver and chloride are 0.75 and 0.76,
respectively.

AgCl(s) = Ag+ + Cl−

We can no longer use the thermodynamic equilibrium constant (i.e. in
absence of diverse ions) and we have to consider the concentration
equilibrium constant or use activities instead of concentration if we
use Kth:

We have calculated the solubility of AgCl in pure water to be 1.0 x
10−5 M, if we compare this value to that obtained in presence of
diverse ions we see % increase in solubility = (1.3 x 10−5 –
1.0 x 10−5) / 1.0 x 10−5 x 100 = 30% Therefore, once again we
have an evidence for an increase in dissociation or a shift of
equilibrium to right in presence of diverse ions.
References[edit]